We can explain why the bond angle of $\ce{NF3}$ (102°29') is lesser than $\ce{NH3}$ (107°48') by the VSEPR theory, since lone pair lone pair repulsion is greater than lone pair bond pair repulsion. Then for $\ce{PH3}$ and $\ce{PF3}$, also, it is expected that the bond angle of $\ce{PF3}$ will be smaller. But I found that the thing is just reverse. Why is this so? Is it an exception?


VSEPR works by accident in the cases of $\ce{NH3}$ and $\ce{NF3}$. In reality, there are much more things to consider as shown in this answer. All four compounds should have bond angles of $90^\circ$ if there were no other effects present.

For both nitrogen compounds, the effects are the short $\ce{N-X}$ bonds which lead to steric clash of the three substituents. Thus, the bond angle is broadened from the original, hypothetical $90^\circ$ to whereever it may seem comfortable to the molecule. Since fluorines are larger atoms than hydrogens are, the $\ce{N-F}$ bond is longer than the $\ce{N-H}$ bond ($137~\mathrm{pm}$ versus $102~\mathrm{pm}$) and therefore the atoms can agree on a smaller bond angle (less crowding).

For phosphorous, we can initially assume that the bond angle of $\ce{PH3}$ is close enough to $90^\circ$ for no hybridisation to be necessary (semi-proven by this answer of Martin) and likewise for $\ce{PF3}$. The bond lengths are $142~\mathrm{pm}$ and $156~\mathrm{pm}$, respectively. The difference in bond lengths is only half of that of the nitrogen compounds ($14$ versus $35~\mathrm{pm}$). Therefore, we can assume the bond length to be mainly dictated by the central phosphorous atom, while the bond angles must be dictated more by the outer atoms. Since again fluorine is larger than hydrogen, this time the better way to stabilise the molecule is not to lengthen the bonds but to expand the angle.


The answer can not be determined using VSEPR theory. The number of lone pairs in both compounds is the same. The answer lies in electronegativity.

  1. In $\ce{NH3}$ nitrogen is more electronegative than hydrogen and therefore it will pull the electrons towards it. Therefore the electrons will be closer to each other hence more repulsion.

  2. But in $\ce{NF3}$ fluorine will pull electrons towards itself therefore the electrons are further comparatively and hence repulsion will be less and so is the bond angle.

But in case of $\ce{PF3}$ and $\ce{PH3}$ back bonding is possible in $\ce{PF3}$ therefore it has a larger bond angle .

  • 1
    $\begingroup$ What type of backbonding are you envisaging? $\endgroup$
    – bon
    Aug 21 '15 at 20:21
  • $\begingroup$ Fluorine is also larger than hydrogen so sterics have a contribution. $\endgroup$
    – gsurfer04
    Aug 21 '15 at 20:36
  • $\begingroup$ The Backbonding is PpiDpi .! $\endgroup$ Aug 22 '15 at 3:56
  • 1
    $\begingroup$ Actually it forms a partial double bond between P and all the three flourine . $\endgroup$ Aug 25 '15 at 19:54
  • 5
    $\begingroup$ Phosphorus does not access d-orbitals. $-1$ $\endgroup$
    – Jan
    Nov 11 '15 at 14:53

In PH3 and PF3 bond angle of PF3 is greater as in PF3 back bonding takes place. Back bonding is possible in PF3 as P has vacant d orbital(as its atomic no. is 15 therefor its electronic configuration is 1s(2e) 2s(2e) 2p(6e) 3s(2e) 3p(3e) this shows it has vacant 3d orbital which can perform back bonding with fluorine as fluorine has pair of electrons. as backbonding takes place the lonepair of P will be used in doublebond formation with F. and obviously if their is no lonepair then their is no one to repell. there for more bond angle also in PH3 no backbonding can take place as hydrogen has no pair ofelectron(e) for backbonding to take place. May this solution help you.:)

  • $\begingroup$ That's not true - significant usage of d-orbitals in p-block element's compounds is outdated theory $\endgroup$
    – Mithoron
    Jan 8 '16 at 13:07
  • $\begingroup$ lol this is the logic outdated theory all concepts of chemistry are years old and you creating your own theories I would say go for it $\endgroup$
    – vedant
    Jan 8 '16 at 13:19
  • $\begingroup$ Chemistry marches on see for ex. chemistry.stackexchange.com/a/18544/9961 for newer description, as this was experimentally disproved. $\endgroup$
    – Mithoron
    Jan 8 '16 at 15:17

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